Remarkable Inertness of Copper(II) Chelates of Cyclen-Based

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Remarkable Inertness of Copper(II) Chelates of Cyclen-Based Macrobicycles with Two trans-N‑Acetate Arms Catarina V. Esteves,† Pedro Lamosa,† Rita Delgado,*,† Judite Costa,‡ Pauline Désogère,§ Yoann Rousselin,§ Christine Goze,§ and Franck Denat*,§ †

Instituto de Tecnologia Química e Biológica, Universidade Nova de Lisboa, Av. da República, 2780-157 Oeiras, Portugal iMed.UL, Faculdade de Farmácia, Universidade de Lisboa, Av. Prof. Gama Pinto, 1649-003 Lisboa, Portugal § Institut de Chimie Moléculaire de l’Université de Bourgogne, UMR CNRS 6302, Université de Bourgogne, 9, Av. Alain Savary, 21078 Dijon, France ‡

S Supporting Information *

ABSTRACT: Two cross-bridged cyclen-based macrocycles with two trans-N-acetic acid arms, one having a dibenzofuran (DBF) moiety as the bridge, H2L1, and the other a diphenyl ether (DPE) one, H2L2, were synthesized. Both compounds behave as “proton sponges.” The thermodynamic stability constants for the Cu2+, Zn2+, Al3+, and Ga3+ complexes of both compounds were determined. They exhibit an excellent thermodynamic selectivity for copper(II), ensuring that metal ions largely present in the human body do not interfere with the copper(II) chelates. All complexes are very slow to form, and [CuL2] and [CuL1] are extremely inert to demetallate, especially [CuL2]. The acid-assisted dissociation of [CuL1] led to a half-life of 4.28 h in 5 M HCl at 363.2 K, while [CuL2] needed harsher conditions of 12 M HCl at 363.2 K with a half-life of 30.8 days. To the best of our knowledge, [CuL2] exhibits the highest half-life value for a copper(II) complex of a polyazamacrocycle derivative reported in the literature until now. Single crystal X-ray diffraction determined for [Cu(H2L1)](ClO4)2 showed the copper center in a distorted octahedral environment bound to the N4O donors of the macrobicycle and one oxygen atom from a carboxylic arm, while for [CuL2] it showed the copper center in a trigonal bipyramid geometry only bound to the donors of the macrobicycle and leaving the carboxylate arms away from the coordination sphere. UV−vis−NIR and X-band EPR spectra showed that in [CuL1] the copper center adopts a distorted compressed octahedral environment, which is the only structure found in solution for this complex, while in [CuL2] a similar environment was found in the first stages of its slow formation but reached a square-pyramidal geometry upon stabilization. The acetate arms play therefore an important role during the formation of the complex, as revealed by the comparison of its complexation behavior with the corresponding parent compounds.

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INTRODUCTION

medicine for radioimmunotherapy using several radiolanthanide chelates,13−16 upon conjugation to suitable biomolecules for target-specific delivery to organs or tissues in patients.17 In order to avoid metal transchelation into the living body, more rigid compounds derived from cyclen or cyclam have been investigated, currently known as constrained or reinforced macrocycles. In these compounds an alkylene bridge is covalently bound to two opposite nitrogen atoms, ethylene or propylene chains are the most common, and the compounds are known as cross-bridged compounds, adamanzanes, or cryptands.18−26 The cross-bridged compounds adopt conformations where all four nitrogen lone pairs point to inside the cavity upon protonation or metal chelation. The small cavity generated and the number of donor atoms particularly

Tetraazamacrocycles and their many derivatives have an excellent ability to form stable complexes with a large scope of metal ions; therefore their study continues to raise huge interest among researchers.1−6 The high thermodynamic stability and kinetic inertness of metal complexes of some Nfunctionalized derivatives of tetraazamacrocycles, especially cyclen (cyclen = 1,4,7,10-tetraazacyclododecane) and cyclam (cyclam = 1,4,8,11-tetraazacyclotetradecane), make such chelates useful for medical applications in clinical imaging and therapy. A large number of cyclen and cyclam derivatives have been prepared with functionalized substituents such as carboxylic acids, phosphonic acids, phosphinic acids, amines, amides, alcohols, pyridyl, and more recently a combination of different substituents in the same molecule.1,2,4−7 Some of them were explored as contrast agents for magnetic resonance imaging (MRI) using Gd3+ complexes8−12 and others in nuclear © 2013 American Chemical Society

Received: January 7, 2013 Published: April 11, 2013 5138

dx.doi.org/10.1021/ic400015v | Inorg. Chem. 2013, 52, 5138−5153

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favors the coordination of the Cu2+ metal ion, especially when two additional arms are bound to the other two trans-N amines, such as acetate, methylphosphonate, methylamide, etc. In general, they are remarkably efficient “proton sponges.”18−26 Due to the geometric configuration of these compounds, their copper(II) complexes exhibit slow formation and inertness to dissociation. These features place these copper(II) chelates as the best candidates for radiopharmaceutics using 67Cu for therapeutic purposes or 64Cu for positron emission tomography (PET).27−34 The main drawback of these chelators for the preparation of radiochelates is their slow kinetic formation, thus needing heating which is not compatible with labeling of sensitive biomolecules (antibodies, enzymes, etc.). In this context we decided to prepare and study two crossbridged derivatives of cyclen, containing unusual and longer bridges bound to the 1,7-positions of cyclen, such as a dibenzofuran, DBF, and diphenylether, DPE, moieties (H2L1 and H2L2, see Chart 1). Both compounds have additional N-

as their coordination properties with Cu2+, Zn2+, Al3+, and Ga3+ ions in solution and in the solid state.

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RESULTS AND DISCUSSION Syntheses of the Compounds. The parent macrobicyclic compounds were synthesized before,35,36 and the additional trans-acetate arms were introduced by reaction with tert-butyl bromoacetate at 40 °C in CH3CN followed by deprotection in concentrated hydrochloric acid, see Scheme 1. Scheme 1. Synthetic Proceduresa

Chart 1. Structures of Compounds Discussed in This Work

a

Reagents and conditions: (i) 2.8 equiv. NaBH(OAc)3, dichloroethane, N2, room temperature, 2 d; (ii) 2.2 equiv. tbutylbromoacetate, 4 equiv. K2CO3, CH3CN, 40 °C, 4 h; (iii) HCl, room temperature, 10 min.

Acid−Base Behavior of the Cross-Bridged Compounds, H2L1 and H2L2. Potentiometric Measurements. Both compounds have six basic centers, four amines and two carboxylates, from which only three protonation constants were possible to obtain by potentiometric data in aqueous solutions, where reliable pH values are only available at the 2.5−11.5 range. The first protonation value was determined by spectrophotometric titrations. 1H NMR titrations were also undertaken at 6.5−14.5 pD values aiming the determination of K1H but without success, see below. All the potentiometric and spectrophotometric measurements were carried out at 298.2 ± 0.1 K in aqueous solution and at an ionic strength of 0.10 ± 0.01 M in N(CH3)4NO3. The stepwise protonation constants (KiH) obtained are compiled in Table 1, together with the values for the parent macrocycles,36 cyclen,37−40 H2do2a,41 and H4dota42,43 for comparison. Overall protonation constants

acetate arms bound to the 4,10-positions of the macrocycle. The parent macrobicycles (L3 and L4) were already studied by some of us,35 including some of their metal complexes.36 They also behave as “proton sponges,” and their complexes with Cu2+ and Zn2+ were very difficult to form and could not be prepared in aqueous solution but in acetonitrile. The introduction of the two acetate arms changes dramatically their metal complexation behavior and improves their solubility in water. In this work, the acid−base reactions of H2L1 and H2L2 were studied as well 5139


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Table 1. Stepwise Protonation (KiH) Constants of H2L1 and H2L2 Compounds, the Parent Cryptands, Cyclen, H2do2a, and H4dota (T = 298.2 K, I = 0.10 M in N(CH3)4NO3)

a

equilibrium reaction

H2L1

H2L2

L3

L4

cyclen

H2do2a

H4dota

L + H+ ⇄ HL HL + H+ ⇄ H2L H2L + H+ ⇄H3L H3L + H+ ⇄ H4L H2L + 2 H+ ⇄ H4L L + 4 H+ ⇄ H4L

15.0a 9.59f 2.33f 1.7f 4.0f 28.6

14.6a 8.86f

>12b 8.97b 2.70b

11.53b 9.05b 2.63b

23.21

11.45d 9.54d 4.00d 2.36d 6.36 27.35

12.09e 9.76e 4.56e 4.09e

>23.7

11.27c 9.96c 2.18c 1.74c 3.92 25.15

5.05g 28.5 b

36 c

30.50 37

This work, determined by spectrophotometric titration. T = 298.2 K, I = 0.1 M in N(CH3)4Cl. T = 298.2 K, I = 0.5 M in KNO3. Slightly different values were found by other authors: T = 298.2 K, I = 0.1 M in NaNO3, log K1 = 10.6, log K2 = 9.6, log K3 < 2, log K4 < 2;38 T = 298.2 K, I = 0.10 M in NaClO4, log K1 = 11.04, log K2 = 9.86, log K3 < 2, log K4 < 2;39 and T = 298.2 K, I = 1 M in NaCl, log K1 = 10.7, log K2 = 9.7, log K3 = 1.73, log K4 = 0.94.40 dT = 298.2 K, I = 0.10 M in N(CH2CH3)4NO3.41 eT = 298.2 K, I = 0.10 M in N(CH3)4NO3.42,43 fThis work, determined by potentiometric titration. gThe values of log K3 and log K4 are on the same order; the system converges with better statistic parameters for a value of the global constant (including K3 and K4) of 5.05, which is about 2.5 for each stepwise constant.

Figure 1. Species distribution diagrams of H2L1 (left) and H2L2 (right) in aqueous solution at cL = 1.0 × 10−3 M.

(βiH) with standard deviations are collected in Table S1 in the Supporting Information. In Figure 1 are represented the species distribution diagrams of H2L1 and H2L2 as a function of the pH.44 These diagrams show the very similar behavior of both compounds along the pH, the neutral species H2L being the main one from pH 3 to about 8, and the completely deprotonated species L2− does not start to form even at pH 12. Depending of the functional group of the arms, all Nsubstituted cyclen derivatives possess two high protonation constants and the others low or very low. Moreover, most of the cross-bridged compounds behave as “proton sponges,” having the first protonation constant particularly high.45 The first two values correspond to the protonation of nitrogen atoms inside the macrocyclic cavity in opposite positions, eventually forming +N−H···N hydrogen bonds.36,45,46 For cyclen, the other two values correspond to the protonation of the other two amines of the ring, and the resulting strong repulsions led to the very small values. In N-acetate cyclen derivatives, when the protonation occurs in the arms bound to nonprotonated amines, the log values are >4, as in acetic acid (and also found for the H4dota).42,43 On the other hand, when the protonation occurs at the acetate arm bound to a protonated amine the value is much lower (≤2.5, as for amino acids) and can occur simultaneously with partial protonation of the last two amines of the macrocycle.41,43,47 On the basis of these data, it is possible to predict that the first two protonations in H2L1 and H2L2 occur mainly on the amines carrying the acetate arms and the next two on the acetate arms, see 1H NMR titrations below.

Spectrophotometric Measurements. The present compounds, as other cross-bridged ones, were found to be “proton sponges.”45 Macrobicycles of short bridges (ethylene and propylene) when protonated adopt a conformation having the four nitrogen lone pairs convergent to the center of the cavity with the protons involved in a network of strong +N−H···N hydrogen bonds. In such cases, the completely deprotonated species is not observed in aqueous solution, and therefore the first protonation constant (K1H) cannot be determined by usual potentiometric measurements. In this work, the K1H values of both compounds were determined by spectrophotometric titrations, using very concentrated solutions of KOH as the titrant, see Figure 2 for H2L2 and Figure S1 of the Supporting Information for H2L1. In fact, by 1H NMR in D2O no shift of proton resonances was observed at pD > 11, and consequently, it was not possible to determine the K1H but only the K2H values, those that were also possible to determine by the potentiometric data. NMR Measurements. The 1H NMR titrations in D2O showed some important differences in the structural behavior of the two compounds derived from the specific architecture of each bridge. The titration of both compounds was carried out at pD values >6. The assignment of the NMR signals of H2L1 was done at 308.2 K and pD 7.48 and 13.78 (see Table S2 of Supporting Information), on the basis of 1D and 2D spectra (COSY, HSQC, and HMBC). The downfield region of the proton modulated carbon spectrum of this compound (APT) displays seven signals, four of which correspond to quaternary carbon 5140


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The assignment of the 1H and 13C resonances of H2L2 was performed at 313.2 K and pH 7.17 (and at pH 13.89, see Table S2) and followed the same reasoning resorting to the COSY and HSQC spectra and taking into consideration the 2 to 3 bond distance information of the HMBC. In H2L2, the geometry of the bridge probably forces the planes of the benzene rings to tilt toward the cyclen plane in opposite directions. This conformation, while preserving the symmetry of the bridge, distorts the symmetry of the cyclen ring. This causes the protons at position 7 to resonate at very different chemical shifts due to the different ring current shifts arising from the different angles adopted in respect to the planes of the aromatic rings (see Table 2) and also the carbon atoms at positions 8 and 9 on either side of the ring to be distinguishable. Differences between the hydrogen atoms above or below the plane of the cyclen ring are also observable, see Figure 4. For H2L1, when the equilibrium H2L1 ⇄ HL1− + H+ occurs (9.39 to 10.95 pD region), for each amount of KOD added the signals of all proton resonances decrease in intensity, and upfield shifted signals of increasing intensity appear, indicating the slow exchange of the equilibrium compared to the NMR relaxation time, see Figure 3. Additionally, it was observed that with the change of the pH, all protons (of the cyclen ring, the acetate arms, and the bridge) exhibit shifts of comparable magnitude. This feature can have two explanations: (a) the proton involved is shared by the four amine centers or/and (b) the deprotonation causes a conformational change which leads to a different orientation of the planes of the aromatic rings located on the bridge, with the consequent difference of the magnitude of the ring current shifts felt by all of these protons. In this last scenario, the effect of the current shifts will obscure the chemical shift changes caused directly by the deprotonation, making it impossible to say, from the NMR data, which of the nitrogen atoms is being deprotonated. Moreover, with these data it was also possible to determine the value of K2D2O, using the ratio of areas between the decreased first signal and the corresponding increased one of the new species, which is log K2D2O = 9.92, see Figure 5. The conversion of this value to the corresponding constant in water gives log K2H2O = 8.948 and 9.3,49 which are values only slightly lower than the more accurate one determined by potentiometric measurements. For this compound, it was not possible to perceive the

Figure 2. UV titration of H2L2 for the determination of the first two protonation constants, at 298.2 K using 10 selected wavelengths. Increased amount of titrant KOH added from 1 to 10.

atoms. The correlations found in the HSQC spectrum led to the identification of the signals of the protons linked to the other three carbon atoms. The sequential information from the COSY spectrum together with the 2 to 3 bond distance information taken from the HMBC spectrum allowed unequivocal assignment of all of the aromatic resonances (see Table 2). Moreover, the HMBC spectrum showed unique correlations between the aromatic signals and a resonance at 57.0 ppm in the 13C spectrum, allowing the assignment of the aliphatic carbon of the bridge (see Figures S2−S5 in Supporting Information). A similar reasoning allowed the assignment of a broad signal centered at 3.05 ppm in the proton spectrum to the CH2 of the acetate groups. The resonances due to the cyclen ring were assigned also using the HMBC spectrum: there were correlations between the bridge and the carbon atom at position 8 (see numbering scheme in Figure 3) and between the acetate arms and positions 8 and 9. It is curious to notice that there are two different proton resonances at position 9; this is probably due to the hydrogen atoms above or below the plane of the cyclen ring. However, this effect is not visible on the protons at position 8, whose signals appear as a singlet.

Table 2. 1H and 13C Resonance Assignment of Compounds H2L1 (at pD = 7.48) and H2L2 (at pD = 7.17) H2L1/δ (ppm) positiona 1 2 3 4 5 6 7 8 8′ 9 9′ 10 11 a

13

C

153.2 121.7 121.8 124.2 129.2 124.3 57.0 48.6 53.8 50.6 168.6

H2L2/δ (ppm) 1

13

H

8.21 (d, 3J = 7.6 Hz) 7.53 (t, 3J = 7.5 Hz) 7.57 (d; 3J = 7.3 Hz) 4.28 (s) 3.14−3.12 3.95−3.93; 3.47−3.43 3.08−3.01

1

C

H

154.9 120.3 130.4 124.6 133.1 127.3 53.5 47.8 47.2 51.3 51.3 54.7 168.5

7.09 7.50 7.36 7.57

(d, 3J = 8.2 Hz) (t, 3J = 7.5 Hz) (t, 3J = 7.2 Hz) (d, 3J = 7.3 Hz)

5.08−5.04; 3.44−3.37 3.59−3.48; 2.98−2.92 3.19−3.12; 2.88−2.82 4.22−4.15;3.61−3.53 3.61−3.53; 3.42−3.25 3.20−3.17

See Figures 3 and 4 for numbering of the atoms. 5141


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Figure 3. 1H NMR spectra of H2L1 in D2O at different pD values and the structure of the compound with numbering of atoms. The pD values are a, 8.23; b, 8.89; c, 9.39; d, 9.67; e, 9.92; f, 10.20; g, 10.60; h, 10.95; i, 12.23; j, 14.39.

Figure 4. 1H NMR spectra of H2L2 in D2O at different pD values and the structure of the compound with numbering of atoms. The pD values are a, 7.34; b, 8.58; c, 8.84; d, 9.05; e, 9.86; f, 10.58; g, 11.22; h, 12.13; i, 13.14; j, 14.35.

deprotonation start of the last proton, even at pD ≈ 14. This proton should be very well accommodated into this rigid bridged cyclen cavity probably forming bifurcated +N−H···N hydrogen bonds with all the other nitrogen atoms, as the resulting 1H NMR spectrum of this species presents only three resonances at the upfield region and two resonances downfield. For the compound H2L2, some of the proton resonances are difficult to follow with pD, especially for the cyclen aliphatic signals (Figure 4). The protons 7 at 5.0 and 3.4 ppm collapse into a broad signal around 4 ppm. The protons at positions 9 and 10 that are detached from the other cyclen aliphatic protons showed a strong downfield shift, indicating the nitrogen bearing the acetate arm as the likely site of this deprotonation, an observation in accordance with the potentiometric pKa values. For the aromatic signals, no

significant shift with the pD was observed for protons 5 and 4, while protons 3 and 2 move slightly, and in opposite directions, see Figure S6. One must highlight that between 8.58 < pD < 9.86 all aliphatic peaks are broad, which can be related to the existence of the second protonation constant and an intermediate exchange rate between the various differently protonated species in respect to the NMR time. On the basis of shifts of only 3, 7, and 10 proton resonances, the value obtained for the corresponding log K2D2O = 9.86. The conversion of this value to the corresponding constant in water gives log K2H2O = 8.8648 and 9.21,49 which are in good agreement with the one determined by potentiometric measurements. Complexation Studies. The ability to form thermodynamically and kinetically stable metal complexes with H2L1 and H 2 L2 was also investigated, and in some cases the 5142


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> 11. Faster complex formation is possible by preparing the solution at pH values > 7.5. Both ligands also form complexes with Al3+ and Ga3+ cations, at low pH, but they are not sufficiently strong to compete with the formation of the corresponding metal hydroxides, see Figures S8 and S9 of the Supporting Information. However, the [GaL1]+ complex is the only species formed in the 3 to 5 pH region, and between about 5 and 7 pH values both ligands form stable species with Al3+ as AlL+ or hydroxocomplexes. When compared to nonbridged cyclen N-acetate derivatives, the stability constant values for the complexes of Cu2+, Al3+, and Ga3+ with the chelators studied here are very high, except for the Zn2+ ions (Table 3). However, stability constants are not the correct way to evaluate the metal complexation ability of ligands exhibiting different basicity. To take this effect into account, the pM values (−log [M]) were calculated at pH 7.4 (physiological) and presented in Table 4. The same information can be retrieved from apparent (or effective) constants determined at different pH values, see Figure S10. Both chelators are selective for copper(II), while H4dota, although exhibiting a slightly higher pCu value, is not selective for this metal ion as it forms complexes with similar strength with all the first transition divalent metal ions including Zn2+, see Figure 7 and Figure S11 for H2L2 and H2do2a.42,43 For Al3+ and Ga3+, the three ligands present very similar pM values, and the distribution curves (Figures S8 and S9 of Supporting Information) indicate that at pH 7.4 the [GaL1]+ and [GaL2]+ complexes are almost completely dissociated, leaving in solution only gallium(III) hydroxides. For Al3+, the studied chelators are slightly better than H4dota existing at physiological pH in the form of hydroxocomplexes. In conclusion, H2L1 and H2L2 are chelators exceptionally selective for Cu2+, coordinating the other divalent metal ions (Mg2+, Ca2+, Ni2+, Zn2+, Cd2+, and Pb2+) with very low thermodynamic stability. For the trivalent Al3+ and Ga3+ cations, the stability constants with H2L1 and H2L2 are larger than for H4dota (Tables 3 and 4) but not enough for it to be possible to use them at physiological pH in medical applications due to the strong tendency of these ions to form hydroxides, see Table 4. Another feature to stress is the crucial role on metal complexation of the acetate arms in H2L1 and H2L2. In fact, the parent compounds, L3 and L4, cannot form complexes

Figure 5. Molar fraction of H2L1 as a function of pD.

thermodynamic stability constants were determined in aqueous solution, at T = 298.2 ± 0.1 K and 0.10 M ionic strength in N(CH3)4NO3. Under these conditions, the kinetics of complex formation is very slow, and as a consequence the determinations were only possible using batch titrations after at least 15 days of stabilization. Once completely formed, the complexes are very inert to demetalation, see below. The obtained values are compiled in Table 3 and Table S3 of the Supporting Information together with those of related ligands for comparison.38,41−43,50,51 The corresponding species distribution diagrams are shown in Figure 6 and in Figures S7−S9 for complexes of Cu2+, Zn2+, Al3+, and Ga3+, respectively. Among the metal ions studied, only Cu2+, Al3+, and Ga3+ can form complexes under the mild conditions used, while the Zn2+ complex ([ZnL]) is formed in small amounts coexisting with Zn(OH)2, see Figure S7 of the Supporting Information. The Ni2+, Cd2+, Pb2+, and Gd3+ form complexes with both ligands in minor amounts, indicating that under the mild conditions used they could not be formed. However, only with Cu2+ are the two chelators able to form complexes in the entire pH range (see distribution diagrams in Figure 6), coexisting as a monoprotonated species [CuHL]+ with free copper(II) at pH < 4, as [CuL] in the 6.0 to 7.5 pH region, and then the [CuL(OH)]− starts to be formed, which is the only species in solution at pH

Table 3. Stepwise Stability Constants (KMHiL), in log units, for the Metal Complexes of H2L1 and H2L2 with Cu2+, Zn2+, Al3+, and Ga3+ Metal Ions in Aqueous Solution (T = 298.2 K and I = 0.10 M in N(CH3)4NO3) equilibrium reactiona

H2L1b

H2L2b

CuL + H+ ⇄ CuHL Cu2+ + L ⇄ CuL CuLOH + H+ ⇄ CuL ZnL + H+ ⇄ ZnHL Zn2+ + L ⇄ ZnL ZnLOH + H+ ⇄ ZnL AlL + H+ ⇄ AlHL Al3+ + L ⇄ AlL AlLOH + H+ ⇄ AlL AlL(OH)2 + H+ ⇄ AlLOH GaL + H+ ⇄ GaHL Ga3+ + L ⇄ GaL

4.39 23.25 9.52

3.85 20.87 9.4

12.4 9.5 3.47 20.17 4.82 8.3

11.59 8.4 4.5 18.3 4.3 8.5

25.08

21.6

cyclenc

H2do2ad

H4dota

23.4

3.0 21.1

3.78e 22.25e

16.2

4.0 18.2

4.18e 21.10e

17.0f

4.00g 21.33g

a

Charges in the complex species were omitted, because they are different for the various ligands. bThis work. cT = 298.2 K, I = 0.1 M in NaNO3.38 T = 298.2 K, I = 0.1 M in N(CH2CH3)4NO3.41 eT = 298.2 K, I = 0.1 M in N(CH3)4NO3.43 fT = 298.2 K, I = 0.20 M in NaNO3.50 gT = 298.2 K, I = 0.1 M in KCl.51 d

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Figure 6. Species distribution diagrams for the complexes of Cu2+ with H2L1 and H2L2 in aqueous solution, cL = cM = 1.0 × 10−3 M.

Table 4. The pMa Values Calculated for Metal Complexes of the Discussed Ligands at pH 7.4 ion 2+

pCu pZn2+ pAl3+ pGa3+

H2L1b

H2L2b

cyclenc

H2do2ad

H4dota

13.19 5.05 12.77 17.83

12.11 5.06 12.79 17.83

18.00 10.80

14.91 12.01

15.19e 14.04e 11.17f 17.83g

octahedral (Oh) environment bound to the four nitrogen atoms of the macrocycle, one oxygen atom from a carboxylic pendant arm and one oxygen from the dibenzofuran (DBF) unit. The octahedron is delimited by an equatorial plane, defined by the two nitrogen donor atoms [N(1), N(3)] and the two oxygen donor atoms [(O3), O(5)], and the compressed axial axis is defined by the nitrogen donor atoms [N(2), N(4)]. The short axial distance is in accordance with the EPR spectroscopic study (see below). The distances between the copper center and N(2), N(3), and N(4) nitrogen atoms, ranging from 2.005(5) to 2.089(2) Å, are typical of such interatomic distances reported in the International Tables for Crystallography [2.071(93) Å]53 or in the literature. The Cu−O(5) distance of 2.116(2) Å is shorter than the ones found in the literature for Cu−O(DBF) bonds, which is of 2.387(35) Å.54,55 This is probably imposed by the strain of the ligand. In fact, a similar distance of 2.131(5) Å was reported for the structure of [CuL3](PF6)2·CH3CN, recently published by Bazzicalupi et al.36 The N−Cu−N and N−Cu−O(5) angles are close to the right angle value expected [86.11(2) to 93.7(8)°]. The distances C(10)−O(1) of 1.197(3) Å and C(12)−O(3) of 1.214(3) Å relative to the distances C(10)−O(2) of 1.332(3) Å and C(12)−O(4) of 1.320(3) Å confirm the carboxylic acid form of the COOH group.

a

Values calculated for 100% molar excess of the ligand over the metal ion with cM = 1.00 × 10−5 M, based on the protonation and stability constants of Tables S1 and S3 or in literature ones. bThis work. cRef 38. dRef 41. eRefs 42 and 43. fRef 50. gRef 51.

with Cu2+ and Zn2+ at room temperature and aqueous solution as found by Bencini et al.36 and by us in this work. Structural Investigations. Single crystals of the compound [Cu(H2L1)](ClO4)2 were obtained from slow evaporation of an aqueous solution of the complex in 5 HClO4 M. The molecular structure is built up from a symmetric unit composed of one [Cu(H2L1)]2+ and two disordered ClO4− as counterions. The ORTEP view52 of the structure of the [Cu(H2L1)]2+ complex cation with the labeling scheme adopted is depicted in Figure 8. In Table 5 are given selected bond lengths and angles for the [Cu(H2L1)]2+ complex cation. The compound acts as a hexadentate ligand. The copper(II) center is in a distorted

Figure 7. Species distribution diagrams for mixtures of Cu2+, Zn2+, and H2L1 (left), and the same cations and H4dota (right) in a 1:1:1 molar ratio and at a concentration of 1.0 × 10−3 M. The amounts of free ligands and metal complexes are expressed as the sum of all of their protonated species in percentage relative to the initial amount of H2L1 or H4dota. 5144


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Figure 8. ORTEP view of [Cu(H2L1)]2+ showing thermal ellipsoids at the 50% probability level. Hydrogen atoms on carbon atoms and counterions have been omitted for clarity.

Table 5. Selected Bond Lengths (Å) and Angles (deg) of [Cu(H2L1)]2+ Complex Cation bond lengths/Å Cu−N(1) Cu−N(2) Cu−N(3) Cu−N(4) Cu−O(5) Cu−O(3) C(10)−O(1) C(10)−O(2) C(12)−O(3) C(12)−O(4)

2.183(2) 2.005(2) 2.089(2) 2.010(2) 2.116(2) 2.557(2) 1.197(3) 1.332(3) 1.214(3) 1.320(3)

bond angles/deg N(1)−Cu−N(2) N(1)−Cu−N(3) N(1)−Cu−N(4) N(1)−Cu−O(5) N(2)−Cu−N(3) N(2)−Cu−N(4) N(2)−Cu−O(3) N(2)−Cu−O(5) N(3)−Cu−N(4) N(3)−Cu−O(3) N(3)−Cu−O(5) N(4)−Cu−O(3) N(4)−Cu−O(5) O(3)−Cu−O(5)

86.36(8) 120.4(9) 86.11(9) 93.70(8) 89.46(9) 170.2(9) 91.0(9) 93.45(8) 88.95(9) 73.58(8) 145.90(8) 97.82(1) 93.28(8) 72.4(7)

Figure 9. ORTEP view of [CuL2] showing thermal ellipsoids at the 50% probability level. Hydrogen atoms on carbon atoms and water molecules have been omitted for clarity.

Further details on the hydrogen bonds and crystal packing are available in the Supporting Information, Table S4 and Figures S12 and S13. The molecular structure of the copper(II) complex of H2L2 is built up from a symmetric unit composed of one [CuL2] and water molecules. The ORTEP view52 of the structure of [CuL2] complex with the labeling scheme adopted is depicted in Figure 9, and in the Table 6 are given selected bond lengths and angles. The compound acts as a pentadentate ligand in the solid state. The copper(II) center is in a trigonal bipyramid (D3h) environment bound to the four nitrogen atoms of the macrocycle and one oxygen from the diphenyl ether (DPE) unit. The copper atom shares a plane with O(5), N(4), and N(1) atoms, which are in equatorial positions at about 120° angles to each other, and with N(2) and N(3) atoms above and below the plane, which are in apical positions. In fact, the trigonal distortion was calculated for the complex using the index structural parameter (τ) as previously defined by Addison et al.56 The τ parameter assumes values of 0 and 1 for ideal

Table 6. Selected Bond Lengths (Å) and Angles (deg) of [CuL2] Complex bond lengths/Å Cu−N(1) Cu−N(2) Cu−N(3) Cu−N(4) Cu−O(5) C(25)−O(3) C(25)−O(4) C(5)−O(1) C(5)−O(2)

2.016(2) 2.003(2) 1.993(2) 2.013(2) 2.137(2) 1.259(3) 1.251(3) 1.251(3) 1.250(3)

bond angles/deg N(3)−Cu−N(2) N(4)−Cu−O(5) N(4)−Cu−N(1) N(1)−Cu−O(5) N(2)−Cu−O(5) N(2)−Cu−N(1) N(2)−Cu−N(4) N(3)−Cu−O(5) N(3)−Cu−N(1) N(3)−Cu−N(4)

176.12(8) 119.71(8) 132.27(8) 108.01(7) 91.85(8) 88.53(8) 89.22(8) 91.96(7) 89.66(8) 89.46(8)

square-pyramidal and trigonal-bipyramidal geometries, respectively. The τ value of 0.73 for [CuL2] is entirely consistent with a distorted trigonal bipyramid coordination sphere. 5145


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Figure 10. UV and vis−NIR spectra of [CuL1] (in blue) and [CuL2] (in red) in aqueous solution.

The distances between the copper center and N(1), N(2), N(3), and N(4) nitrogen atoms, ranging from 1.993(2) to 2.016(2) Å, are typical of such interatomic distances reported in the International Tables for Crystallography [2.071(93) Å]53 or in the literature. The Cu−O(5) distance of 2.137(2) Å is shorter than the ones found in the literature for Cu−O(ether) bonds, which is 2.371(96) Å, probably imposed by the strain of the ligand. A similar distance of 2.165(2) Å was reported for the corresponding distance in the structure of [CuL4](ClO4)2.35 Further details on the hydrogen bonds are available in the Supporting Information, Table S5. UV−Vis−NIR and EPR Spectroscopic Studies of the Copper(II) Complexes of H2L1 and H2L2. UV−vis−NIR spectra were performed for [CuL1] and [CuL2] in aqueous solutions at pH 9.27 and 9.16, respectively (up to pH about 7, the spectra did not change). The [CuL1] complex exhibits a broad band in the visible region centered at λ = 650 nm (ε = 86 M−1 cm−1) with a shoulder at higher energy (565 nm, 49 M−1 cm−1), also a broad band in the NIR region (1051 nm, 57 M−1 cm−1) and several bands in the UV region. The [CuL2] complex presents also a large band with λmax = 520 nm (ε = 190 M−1 cm−1) and a shoulder at 668 nm (80 M−1 cm−1), no bands in the NIR region, and in the UV region exhibits only three main bands, see Figure 10 and Table S6. The bands in the vis− NIR region of both complexes, due to the copper d−d transitions, are quite different, indicating dissimilar geometries for the copper centers. The position and intensity of the bands rules out regular octahedral or tetragonal geometries.57−60 However, especially for five and six coordination numbers, it is difficult to infer structural features from electronic spectra of copper(II) complexes as their stereochemistries vary over an appreciable range of distortion within a given coordination number, which is known as the plasticity effect.60 To go further in the characterization of the complexes, Xband EPR spectra of the same solutions were carried out at 90− 298 K, and two of them are presented in Figure 11 together with the simulated ones. The spectra were obtained in a frozen water/ethylene glycol solution (1:1 v/v), although in H2O they are similar, but in better resolution in the mixture of solvents. The EPR parameters obtained by simulation of the spectra of both complexes61 are compiled in Table 7, together with the values for [CuL3]2+ and [CuL4]2+ under the same experimental conditions. For [CuL1] complex, three different values of g were obtained, with gx > gy ≫ gz ≈ 2.00, which is of the “inverted”

Figure 11. X-band EPR spectra of [CuL1] (in blue) and [CuL2] (in red) complexes in a frozen water/ethylene glycol solution (1:1 v/v) recorded at 90 K and the simulated ones (below the corresponding experimental spectra in dots). Experimental conditions: at pH 9.16 and concentration 2.70 × 10−3 M for [CuL1] and pH 9.27 and 2.66 × 10−3 M for [CuL2]. Microwave power of 2.0 mW, modulation amplitude of 1.0 mT, and the frequency (ν) was 9.68 GHz.

type. For these cases, the R parameter can be indicative of the predominance of the dz2 or dx2−y2 orbital in the ground state of the unpaired electron of the Cu2+ ion, R = (gy − gz)/(gx − gy) with gx > gy > gz. When R > 1, the greater contribution to the ground state arises from the dz2 orbital, and when R < 1, the greater contribution to the ground state comes from the dx2−y2 orbital.62,63 The R value of 1.17 determined for [CuL1], as well as the values of the hyperfine splitting parameters, are indicative of a predominance of the dz2 ground state, characteristic of axial-compressed geometries, such as distorted compressed octahedral or trigonal bipyramidal.57−60,62−64 The data suggest the less common distorted compressed octahedral geometry, as also found in the crystal X-ray structure (see above). In spite of the compressed octahedral and trigonal bipyramidal geometries not being able to be distinguished in solution by their EPR spectra, they present very different vis−NIR spectra. In fact, the spectra of copper(II) in trigonal bipyramidal geometry are more intense then those of compressed octahedral one, but the latter present a NIR band almost of the same intensity as that of the vis region, as in the case of [CuL1].58,64 The EPR spectra 5146


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Table 7. EPR Parameters for the Copper(II) Complexes of H2L1 and H2L2 at 90 K in Water/Ethylene Glycol Solution (1:1 v/ v)a EPR parameters complex

pH

[CuL1] [CuL2] [CuL3]2+ [CuL4]2+ [Cu(cyclen)NO3]+d

8.9 9.2 7.9 7.7

λmax (ε)b 647 520 585 501 594

(76.0) (310.7) (15) (69) (271)

gx

gy

gz

Axc

Ayc

Azc

R

2.211 2.03 2.03 2.03

2.117 2.05 2.058 2.06

2.007 2.162 2.176 2.161 2.172

115.9 3.0 2.4 30.2

51.0 24.5 46.8 29.6

75.4 177.0 194.5 204.2 177.5

1.17 0.18 0.24 0.25 0

2.089

31.2

Concentration of the complexes in the range 1.90 × 10−3 to 7.25 × 10−3 M. bλmax of visible absorbance band in nm and εmol in M−1 cm−1. cAi × 104 (cm−1). dValues from refs 69 and 70. a

defined by the remaining secondary nitrogen atoms together with the oxygen of the bridge. In solution, the spectroscopic data revealed the copper center also in a N4O square pyramidal environment for both [CuL3]2+ and [CuL4]2+ complexes, the [CuL3]2+ exhibiting a stronger axial ligand field. Although the spectroscopic data for [CuL2], [CuL3]2+, and [CuL4]2+ point in all cases to square pyramidal environments around the copper centers, the distortions from this geometry are more significant in [CuL2]. In spite of the similarities of geometries, in solution, between [CuL2] and [CuL4]2+, the role of the H2L2 arms in metal chelation is clearly observable. Indeed, L3 and L4 could not form metal complexes,36 including with Cu2+, under aqueous mild conditions. Trying to understand the role of the arms in H2L1 and H2L2 in metal chelation during the slow reactions of these chelators with Cu2+, electronic, and EPR spectra were acquired along time. The results can be observed in Figures S16 and S17 for the vis and X-band EPR spectra, respectively. In the case of H2L1 complexes, the sequence of spectra show the decrease of [Cu(OH2)6]2+ and the increase of [CuL1] amounts, clearly indicating that the species formed since the beginning is the same as the final one, only the amount of complex formed increases with time. However, for the H2L2 complex, a species with different geometry is formed in the first stages. In fact, the vis−NIR showed (see in Figure S18 the vis−NIR spectrum of the intermediate compound in more detail) one band at 630 nm with a shoulder at 520 nm and another larger band centered at 1030 nm. Along time, the band at 520 nm increases in intensity while that at 1030 nm decreases. In the final species, the band at 520 nm is the main one, while the first formed species exists in small percentage. The EPR spectra shown in Figure S17 supports the presence of one species with the copper center in the dz2 ground state together with an increasing amount of the final species in the dx2−y2 ground state. Additionally, the first spectra also revealed the presence of [Cu(OH2)6]2+ that in the beginning is practically the only species in solution. It was not possible to obtain the completely pure complex corresponding to the dz2 ground state, as it exists always with a certain amount of the final species; therefore the simulation of the spectra is not straightforward. However, all the data for this species are very similar to those obtained for [CuL1] (see in more detail one of the spectra with the corresponding simulation in Figure S19). The results clearly indicate that the intermediate copper(II) complex of H2L2 corresponds also to a copper center in a hexacoordinate environment, of compressed octahedral symmetry, although the final stabilized complex point to a distorted square pyramidal geometry. These features suggest that the strain imposed by the H2L2 backbone to the geometry of the copper(II) center is much more significant than that of H2L1, preventing the direct

recorded at several temperatures, starting at room temperature (297 K), showed no significant changes, indicating that the geometry around the copper is preserved, see Figure S14. In contrast, for [CuL2] the values for the g parameter are gz > gy ≈ gx and gx ≥ 2.03, characteristic of copper(II) complexes with slightly rhombic symmetry with elongation of the axial bonds and a dx2−y2 ground state. Elongated rhombic-octahedral, tetragonal, distorted bicapped square pyramidal, or distorted square pyramidal symmetries would be consistent with these data.57−60,64−67 On the other hand, the value of the maximum of the band in the visible and its large intensity point to square pyramidal geometries. The abnormally low value of gz in [CuL2], when compared with the expected values of gz versus Az diagrams of Peisach and Blumberg,68 can be explained by severe distortions of the N4 equatorial plane. Moreover, an increase of the gz and a decrease of Az parameters and simultaneously a red-shift of the d−d absorption band in the electronic spectra with coordination of axial ligands is expected.64 However, the distortions of the Cu geometry in [CuL2] appear to be analogous to those of [Cu(cyclen)(NO3)](NO3),69−71 which adopts a square pyramidal geometry with one oxygen of the NO3− ligand in an axial position and the copper atom 0.5 Å above the plane containing the four N atoms of the macrocycle.71 The observation of identical values for g0 at 297 K (g0 = 2.088 and A0 = 201.9 × 10−4 cm−1) and ⟨g⟩ at 90 K indicates that the solution coordination is not significantly temperature dependent (consistent with g0 = (gx + gy + gz)/3 and A0 = (Ax + Ay + Az)/3, see Figure S15). As seen above, in the single crystal X-ray structure, the copper center is also in a N4O pentacoordinated environment, although adopting a trigonal pyramidal geometry. Data of the electronic and EPR spectra of [CuL3]2+ and [CuL4]2+ under the same experimental conditions are also included in Tables 7 and S6. The complexes of both parent ligands present vis−NIR and EPR spectra in aqueous solution characteristic of copper(II) with elongation of the axial bonds, in a dx2−y2 ground state and distorted N4O square pyramidal geometry. The impact of the strain imposed by the backbone of the ligands is also observed for the complexes of the parent ligands, giving rise to smaller gz values than expected. In the crystal X-ray structures of [CuL3](PF6)2·CH3CN36 and [CuL4](ClO4)2,35 the copper is enclosed into the cavity of the macrobicycle bound to the five donor atoms, in a N4O coordination sphere. In [CuL4]2+, the copper center adopts a square pyramidal geometry where the four nitrogen atoms form the basal plane and the oxygen of the bridge takes the apical position, while in [CuL3]2+ it displays an intermediate geometry between a square pyramid and a trigonal bipyramid one. In the last geometry, the apical positions are occupied by the two tertiary nitrogen atoms, and the equatorial plane is 5147


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most polyazamacrocyclic complexes of copper(II) have rather negative reduction potentials that are well below the estimated −0.40 V (NHE) threshold for typical bioreductants.4 It was observed that cross-bridged cyclam copper(II) complexes typically exhibit quasi-reversible reductions but not those of cross-bridged cyclen ones, suggesting that the former macrobicycle can adapt somewhat to the coordination preferences of Cu(I).72 The cyclic voltammograms of the copper(II) complexes of H2L1 and H2L2 were carried out in aqueous solution at neutral pH, see Figure S21. The reductions of both complexes are irreversible with two cathodic waves at −652 and −784 mV (Ag/AgCl, scan rate 20 mV/s) for [CuL1] and −644 and −788 mV (Ag/AgCl, scan rate 20 mV/s) for [CuL2]. However, only one anodic peak on the reverse sweep was observed at −532 mV for [CuL1] and −584 mV for [CuL2]. This behavior is indicative of coupled chemical events. It seems that the reduced complexes may exist as mixtures of at least two species, which quickly equilibrate. It is curious to observe that the two species have inverse amounts in [CuL1] and [CuL2], see Figure S21b. These species may consist of four- and five-coordinate Cu+ complexes, as also seen for other macrobicyclic complexes.73,74 Therefore, H2L1 and H2L2 ligands, as all cyclen derivatives, due to their rigid backbones and the resulting very small cavity, do not appear to be able to adapt well to the coordination requirements of Cu+ and stabilize it.

coordination of the acetate arms of the ligand in the stable form of the complex. Probably, the copper(II) complex of H2L2 starts to be formed using one (or both) arm(s) on the coordination to the copper in a distorted compressed octahedral geometry, but the distortion to the usual geometries of the copper(II) imposed by the ligand led to a final rearrangement and to a square pyramidal environment, where the carboxylate groups are so distant that they cannot be considered bound to the copper center. Kinetic Stability of [CuL1] and [CuL2] Complexes in Acidic Solution. The kinetic inertness of a complex to the dissociation can be more significant than its thermodynamic stability in the selection of ligands to form radiochelates to be used in medical applications.72 The usual assay for evaluation of the kinetic inertness of a complex is its acid demetalation under pseudo-first-order conditions. It was found that acid inertness half-lives obtained under these conditions are useful first predictors for in vivo viability of 64Cu-labeled chelates.32 Therefore, the kinetic inertness of [CuL1] and [CuL2] was evaluated at 5 and 12 M HCl, respectively, in aqueous solution at several temperatures under pseudo-first-order conditions assayed by monitoring the υmax absorbance in their visible spectra, see Figure S20 for the complex of H2L1. The results are compiled in Table 8, which also includes the values for the Table 8. Half-lives of the Dissociation of [CuL1] and [CuL2] in Very Harsh Conditions, and Values for the Most Kinetically Inert Copper(II) Complexes from the Literature complex [Cu(L1)] [Cu(L2)] [Cu(c3b-te2a)] [Cu(cb-te2a)] [Cu(c3b-do2a)] [Cu(cb-do2a)]

conditions [HCl]/M (T/K) 5 12 12 5 12 5

(363.2) (363.2) (363.2) (363.2) (363.2) (303.2)

half-life (t1/2)

reference

4.28 h 30.8 d −a 6.4 d 1.1 d 2σ(I) and all data, respectively. Crystal Data for [CuL2]. 2(C26H32CuN4O5), 7(H2O), M = 1214.30, monoclinic, space group C2/c, a = 9.6023(4) Å, b = 16.4735(6) Å, c = 33.8442(14) Å, β = 94.425(2)°, V = 5337.6(4) Å3, Z = 4, T = 115(2) K, Dc = 1.511 g cm−3, λ(Mo Kα) = 0.71073, μ(Mo Kα) = 0.878 mm−1, 31 174 reflections collected, 6164 unique. The minimum and maximum residual electron densities are −0.439 and 0.513 e Å−3. The final agreement factors are R(1) = 0.0425 and 0.0789, and wR(2) = 0.0779 and 0.0876, for I > 2σ(I) and all data, respectively. CCDC-907927 and CCDC-924316 contains the supplementary crystallographic data for this paper. These data can be obtained free of charge from The Cambridge Crystallographic Data Centre via www. ccdc.cam.ac.uk/data_request/cif.

sources were used as received. Organic solvents were dried by standard methods.75 Elemental analyses and electrospray mass spectra (ESIMS) were performed by the Analytical Services Unit of ITQB-UNL/ IBET. The 1H and 13C{1H} NMR spectra for the ligand characterization were recorded on two spectrometers, a Bruker Avance III 400 (1H at 400.13 MHz and 13C at 100.61 MHz) and on a Bruker Avance III 800 (1H at 800.33 MHz and 13C at 201.24 MHz), and the pH titration studies were recorded on a Bruker Avance DRX 300 spectrometer (1H at 300.13 MHz) at a probe temperature of 298.2 K. Chemical shifts (δ) are given in parts per million and coupling constants (J) in hertz. The 3-(trimethylsilyl) propionic acid sodium salt was used as an internal reference for 1H spectra in D2O. The resonance assignments are based on peak integration and multiplicity, and on 2D homo- and heteronuclear correlation experiments. Synthesis of Compound L3. A solution of dibenzofuran-4,6dicarbaldehyde (2.88 g, 12.84 mmol) in 1,2-dichloroethane (40 cm−3) was added dropwise to a stirred solution of 1,4,7,10-tetraazacyclododecane (cyclen; 2.21 g, 12.84 mmol) and fresh triacetoxyborohydride (7.61 g, 35.92 mmol) in 1,2-dichloroethane (250 cm−3). The solution was stirred at room temperature under an atmosphere of nitrogen for 48 h. The reaction mixture was quenched by the addition of a 1 M NaOH aqueous solution (300 cm−3), and the product was extracted with chloroform (3 × 200 cm−3). The organic layer was dried over MgSO4 and evaporated under reduced pressure. The crude product was washed with cyclohexane to give the compound L3 as a yellow powder (3.51 g, 75%). 1H NMR (500 MHz, CDCl3, 298 K): δH 2.50 (bs, 2 H), 2.53 (m, 8 H), 2.64 (m, 8 H), 4.00 (s, 4 H), 7.18 (m, 4 H), 7.79 (m, 2 H). 13C NMR (125 MHz, CDCl3, 298 K): δC 47.9, 53.0, 58.6, 120.9, 123.1, 124.4, 125.5, 128.7, 154.8. Elem. Anal. Calcd. for C22H28N4O (%): C, 72.50; H, 7.74; N, 15.37. Found: C, 72.3; H, 7.8; N, 15.3. MALDI-TOF (m/z): 364.9. Synthesis of Compound H2L1. The compound L3 was dissolved in MeCN (350 cm−3), and K2CO3 (3.93 g, 28.52 mmol) and tert-butyl bromoacetate (2.78 g, 14.27 mmol) were then added. The reaction was heated to 40 °C for 4 h. After cooling, solids were removed by filtration, and the filtrate was evaporated under reduced pressure. Residual solvent was removed under high vacuum conditions, and the product was triturated with diethyl ether to give the compound as a yellow oil (2.19 g, 91%). 1H NMR (600 MHz, CDCl3, 323 K): δH 1.26 (s, 18 H), 2.68−2.71 (m, 4 H), 2.88−2.94 (m, 8 H), 3.01−3.05 (m, 4 H), 3.14−3.18 (m, 4 H), 4.67 (s, 4 H), 7.33 (t, 2 H, J = 7.5 Hz), 7.51 (d, 2 H, J = 7.5 Hz), 7.92 (d, 2 H, J = 7.5 Hz), 8.52 (br s, N+−H). 13C NMR (150 MHz, CDCl3, 323 K): δC 27.9, 51.7, 51.8, 53.1, 57.8, 81.9, 119.6, 121.5, 123.7, 123.9, 128.9, 153.8, 168.8. ESI+ (m/z) 593.4 (M + H)+; 537.3 (M + H − tBu)+. The ester groups were hydrolyzed in concentrated hydrochloric acid (15 cm−3). The reaction mixture was stirred at room temperature for 10 min. The hydrochloric acid was then evaporated off. The remaining oil was dissolved in acetone, and the desired product slowly precipitated. After removal of trace solvents under high vacuum conditions, L1(HBr)(HCl)·2H2O (1.487 g) was obtained as a white solid in a 67% yield. 1H NMR (600 MHz, D2O, 363 K, pD = 14): δH 2.52−2.56 (m, 5 H), 2.63−2.71 (m, 5 H), 2.98−3.02 (m, 5 H), 2.99− 3.07 (m, 4 H), 3.12−3.17 (s, 4 H), 3.74 (s, 4 H), 7.31−7.32 (m, 2 H), 7.34−7.36 (m, 2 H), 7.90−7.92 (m, 2 H). Elem. Anal. Calcd. for C26H37BrClN4O7 (%): C, 49.26; H, 6.04; N, 8.84. Found: C, 49.60; H, 6.10; N, 8.93: ESI+ (m/z) 481.2 (M + H)+; 503.2 (M + Na)+. Synthesis of Compound L4. The compound L4 was obtained from cyclen (0.78 g, 4.53 mmol) and bis(2-formylphenyl)ether (1.02 g, 4.53 mmol) in the presence of triacetoxyborohydride (2.69 g, 12.68 mmol) by using the same procedure reported for L3. The compound was isolated as a white powder (1.29 g, 78%). 1H NMR (500 MHz, CDCl3, 298 K): δH 2.05 (bs, 2 H), 2.45−2.82 (m, 16 H), 3.33 (d, 2 H, J = 12.5 Hz), 4.13 (d, 2 H, J = 12.5 Hz), 6.65 (m, 2 H), 6.96 (m, 2 H), 7.15 (m, 4 H). 13C NMR (125 MHz, CDCl3, 298 K): δC = 48.1, 48.4, 53.5, 55.9, 59.2, 119.3, 123.7, 130.0, 130.9, 132.5, 157.3. Elem. Anal. Calcd. for C22H30N4O (MW = 366.51) (%): C, 72.08; H, 8.26; N, 15.29. Found: C, 72.0; H, 8.3; N, 15.2. MALDI-TOF (m/z): 367.1. Synthesis of Compound H2L2. The compound L4 (1.0 g, 2.73 mmol) was dissolved in MeCN (150 cm3). K2CO3 (1.51 g, 10.92 5149


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X-Ray Equipment and Refinement. Diffraction data were collected on a Nonius Kappa Apex-II CCD diffractometer equipped with a nitrogen jet stream low-temperature system (Oxford Cryosystems). The X-ray source was graphite monochromated Mo Kα radiation (λ = 0.71073 Å) from a sealed tube. The lattice parameters were obtained by least-squares fit to the optimized setting angles of the entire set of collected reflections. No significant intensity decay or temperature drift was observed during the data collections. For [Cu(H2L1)]2+, data were reduced by using DENZO software76 without applying absorption corrections; the missing absorption corrections were partially compensated by the data scaling procedure in the data reduction. The structure was solved by the charge flipping algorithm using the SUPERFLIP program.77 For [CuL2], data were reduced by using SAINT V8.27B software78 with applying multiscan absorption corrections. The structure was solved by using the SHELXS program.79 Refinements were carried out by full-matrix least-squares on F2 using the SHELXL97 program79 on the complete set of reflections. Anisotropic thermal parameters were used for non-hydrogen atoms. All H atoms, on carbon or oxygen atoms, were placed at calculated positions using a riding model with C−H = 0.95 Å (aromatic) or 0.99 Å (methylene) or O−H = 0.84 Å with Uiso(H) = 1.2Ueq(CH2), Uiso(H) = 1.2Ueq(CH), or Uiso(H) = 1.5Ueq(OH). In [Cu(H2L1)](ClO4)2, three oxygen atoms of each perchlorate anion exhibited disorder with the ratios 0.59(1)/0.41(1). The geometric parameters of minor disordered components in each group were restrained by using SAME restraints.79 Similar Uij constraints were applied within the disordered parts to maintain a reasonable model by using EADP constraints.79 Potentiometric Equipment and Work Conditions. The potentiometric setup for conventional titrations consisted of a 50 cm3 glass-jacketed titration cell sealed from the atmosphere and connected to a separate glass-jacketed reference electrode cell by a Wilhelm type salt bridge containing 0.10 M N(CH3)4NO3 solution. An Orion 720A+ measuring instrument fitted with a Metrohm 6.0123.100 glass electrode and an Orion 90−05−00 Ag/AgCl reference electrode was used for the measurements. The ionic strength of the experimental solutions was kept at 0.10 ± 0.01 M with N(CH3)4NO3; temperature was controlled at 298.2 ± 0.1 K using Huber CC3-K6 compact cooling and heating bath thermostats and an Orion 91−70−06 ATC-probe previously calibrated. Atmospheric CO2 was excluded from the titration cell during experiments by slightly bubbling purified nitrogen on the experimental solution. Titrant solutions were added through capillary tips at the surface of the experimental solution by a Metrohm Dosimat 665 automatic buret. The titration procedure is automatically controlled by software after the selection of suitable parameters, allowing for long unattended experimental runs. In cases where automatic titrations could not be preformed, out-of-cell titrations were carried out, and the electromotive force was measured with a Metrohm 6.0234.100 combined pH electrode previously calibrated. Potentiometric Measurements. Purified water was obtained from a Millipore Milli-Q demineralization system. Stock solutions of H2L1 and H2L2 were prepared at ca. 2 × 10−3 M. The N(CH3)4NO3 salt was prepared by neutralization of a commercial N(CH3)4OH solution with HNO3. Metal ion solutions were prepared in water at 0.025−0.050 M from analytical grade nitrate salts of the metal ions and standardized by titration with Na2H2edta.80 Carbonate-free solutions of the titrant N(CH3)4OH were obtained at ca. 0.10 M by treating freshly prepared Ag2O with a solution of N(CH3)4I under nitrogen. These solutions were standardized by application of Gran’s method.81 A 0.100 M standard solution of HNO3 prepared from a commercial ampule was used for backtitrations. The [H+] of the solutions was determined by measurement of the electromotive force of the cell, E = Eo′ + Q log[H+] + Ej. The term pH is defined as −log [H+]. Eo′ and Q were determined by titrating a solution of known hydrogen-ion concentration at the same ionic strength in the acid pH region. The liquid-junction potential, Ej, was found to be negligible under the experimental conditions used. The value of Kw = [H+][OH−] was found to be equal to 10−13 80 by titrating a solution of known

hydrogen-ion concentration at the same ionic strength in the alkaline pH region, considering Eo′ and Q valid for the entire pH range. Measurements during conventional titrations were carried out with ca. 0.05 mmol of ligand in a total volume of ca. 30 cm3, in the absence of metal ions and in the presence of each metal ion at 0.9:1 M/L ratio. A backtitration was always performed at the end of each direct complexation titration in order to check if equilibrium was attained throughout the full pH range. Each titration curve typically consisted of 50−60 points in the 2.5−11.5 pH range, and a minimum of two replicate titrations were performed for each system. In all cases, the direct and back-titration curves differ markedly, especially for the copper(II) ones of both ligands. In such cases, out-of-cell titrations were carried out. Out-of-cell titrations (or batch titrations) for all systems of slow kinetics formation were carried out by preparation of independent vials of different pH values and under the experimental conditions used for the conventional titrations at about 3.00 cm−3 total volume. Each vial contains 3.00 cm−3 of the mother solution of the complex with the necessary amount of N(CH3)4NO3 to control the ionic strength, and then the pH was adjusted at the desired value. A mother solution of the complex was prepared by addition of the ligand and metal ion in a 1:1 ratio, and the pH was adjusted at a value >7.5 by the addition of base (N(CH3)4OH) or acid (HNO3) and used after 2−3 h of equilibration. The vials were tightly closed under nitrogen and kept at 298.2 K until the equilibrium was reached, which was controlled each week. For the complexes of both ligands, the equilibrium was generally reached after one week, and the vials were kept under the same conditions one additional week for control reasons. NMR Spectroscopy. Two dimensional spectra were acquired on a Bruker Avance II 500 spectrometer (Bruker, Rheinstetten, Germany) working at a proton operating frequency of 500.43 MHz, equipped with a four channel 5 mm inverse detection probe head with pulsefield gradients along the z axis, or on a Bruker Avance III 800 spectrometer (Bruker, Rheinstetten, Germany) working at a proton operating frequency of 800.33 MHz, also equipped with a four channel 5 mm inverse detection probe head. For the assignment strategy, spectra were run at 308.2 and 313.2 K using standard Bruker pulse programs. 13C spectra were recorded at 125.76 MHz using the APT (attached proton test) sequence. The modulation of peak sign, to distinguish methyl and methyne from methylene signals, was achieved using a delay of 6.89 ms for the evolution of 1JCH. Proton decoupling was applied during the acquisition stage using the WALTZ-16 sequence.82 In the twodimensional 1H−13C heteronuclear single quantum coherence (HSQC) spectra, a delay of 3.45 ms was used for evolution of 1JCH, while in the heteronuclear multiple bond connectivity (HMBC) spectra a delay of 73.5 ms was used for the evolution of long-range couplings. In the HSQC, proton decoupling was achieved using the GARP4 sequence.82 For determination of the first two protonation constants of both ligands, 1H NMR spectra in D2O solution, at 6.5−14.5 pD range and 298.2 K, were recorded, ca. 20 points per titration. The ligand stock solutions were prepared at 0.010 M, and the titrant was a fresh CO2free KOD solution. The titrations were performed directly in the NMR tube, and the titrant was added with a research syringe (0.1−2.5 μL) and/or automatic pipettes (0−25 and 0−100 μL). The pH* was measured with an Orion 420A instrument fitted with a Hamilton Spintrode PN23819703 combined microelectrode after calibration with two buffers prepared in aqueous solution (pH 8.00 of borate/ hydrochloric acid and pH 4.00 of citric acid/sodium hydroxide/ sodium chloride). The final pD was calculated according to the equation pD = pH* + (0.40 ± 0.02),48 where pH* corresponds to the reading of the pH meter. The measurements were carried out with ca. 0.05 mmol of each ligand in a total volume of 0.5 cm3 without control of the ionic strength. The equilibrium constants in D2O (KD) were converted in H2O (KH) values using published equations.48,49 UV−Vis−NIR Measurements. Absorption spectra of the solutions studied were recorded from 200 to 900 nm at T = 298.2 ± 0.1 K using a UNICAM model UV-4 spectrophotometer, and from 850 to 1200 5150


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nm using a Shimadzu model UV-3100 spectrophotometer for NIR ranges. For determination of the first protonation constant of both ligands, UV spectra of solutions of pH > 11.5 were recorded. The solutions were obtained by the addition of known amounts of KOH at 298.2 K, ca. 15 points per titration. Ligand stock solutions (7.0 × 10−5 M) and fresh CO2-free KOH solutions as the titrant were prepared. The titration was carried out directly in the UV cell (with ca. 0.2 μmol of each ligand in a total volume of 3.0 cm3, without control of the ionic strength), and the titrant was added with a Crison microBU 2031 buret. Calculation of Thermodynamic Equilibrium Constants. The data from potentiometric titrations were used to determine the protonation constants of H2L1 and H2L2 (except the first one) and the stability constants with the different metal ions, while the 1H NMR and UV-spectrophotometric titrations in the >6 pH region were used only to calculate the first two protonation constants of both compounds. The overall equilibrium constants βiH and βMmHhL l (being βMmHhLl = [MmHhLl]/[M]m[H]h[L]l and βMH−1L = βML(OH) × Kw) were obtained by refinement of the potentiometric and spectrophotometric data with the HYPERQUAD program83 and of the 1H NMR data by the HYPNMR program.84 Differences, in log units, between the values of protonated (or hydrolyzed) and nonprotonated constants provide the stepwise (log K) constants (being KMmHhLl = [MmHhLl]/[MmHh−1Ll][H]). The errors quoted are the standard deviations of the overall stability constants calculated by the program when all the experimental data (at least two titration curves) for each system were fitted together. Species distribution diagrams were plotted from the calculated constants with the HYSS program.44 Kinetic Measurements. The dissociation kinetics of complexes of H2L1 and H2L2 with the Cu2+ ion in aqueous solution was followed as a function of time by absorption spectroscopy in the 400−850 nm region, using a UNICAM UV−vis spectrophotometer model UV-4, at 298.2 and 363.2 K. The spectra of [CuL1] were recorded each 15 min, and those of [CuL2] each 24 h until the complete disappearance of the band in the visible region. The solutions of the complexes were prepared, using CuCl2 or Cu(ClO4)2, at 5.0 × 10−4 M in 5 M HCl (or 5 M HClO4). The results were interpreted under pseudo-first-order conditions, and the half-life was calculated from the slope of linear ln(absorbance) vs time plots. X-Band EPR Spectra. The EPR spectroscopic measurements were recorded with a Bruker EMX-8/2.7 spectrometer equipped with continuous-flow cryostat for liquid nitrogen operating at X-band. Room temperature (297 K) spectra were recorded on as-prepared solutions injected into a 1.6 mm inner diameter clear fused quartz cell to a height of approximately 10 mm. Solutions of the copper(II) complexes were prepared at 5.0 × 10−3 M. To these solutions ethylene glycol was added for the final 1:1 v/v H2O/ethylene glycol ratio. The EPR spectra were recorded at a microwave power of 2.0 mW, frequency (ν) 9.67 GHz, T = 90−200 K and at room temperature, and were simulated using SpinCount software.61 Electrochemical Studies. A BAS CV-50W Voltammetric Analyzer connected to BAS/windows data acquisition software was used. Cyclic voltammetric experiments were performed in a glass cell MF-1082 from BAS in a C-2 cell enclosed in a Faraday cage, at room temperature, under nitrogen. The reference electrode was Ag/AgCl (MF-2052 from BAS) filled with 3 M NaCl in water, standardized for the redox couple Fe(CN)63−/Fe(CN)64−. The auxiliary electrode was a 7.5 cm platinum wire (MW-1032 from BAS) with a gold-plated connector. The working electrode was a glassy carbon (MF-2012 from BAS). Copper(II) complexes of H2L1 and H2L2 (≈ 6 × 10−3 M; pH ≈ 7) were prepared in 0.10 M N(CH3)4NO3 in water. The solutions were deaerated by a nitrogen stream prior to all measurements and were kept under nitrogen during the measurements. Between each scan, the working electrode was electrocleaned by multicycle scanning in the supporting electrolyte solution, polished on alumina 1 and 0.05 μm,

cleaned with water, and sonicated before use, according to standard procedures. Cyclic voltammograms with a sweep rate ranging from 10 to 200 mV s−1 were recorded in the region from +1.2 to −1.2 V. At this potential range the ligands were found to be electroinactive. The halfwave potentials, E1/2, were obtained by averaging the anodic and cathodic peak potentials. All potential values are reported relative to the Ag/AgCl reference electrode, and the E1/2 and ΔEp of the Fe(CN)63‑/Fe(CN)64− couple, under our experimental conditions, were 196 mV and 73 mV, respectively.

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ASSOCIATED CONTENT

S Supporting Information *

Tables of overall protonation and stability constants, of the assignment of 1H and 13C NMR resonances of H2L1 and H2L2 at basic pH; of UV−vis−NIR data of copper(II) complexes; HSQC and HMBC spectra of H2L1 and H2L2; representative speciation diagrams in solution; diagram of log Keff as a function of the pH for several ligands; UV−vis spectra of the Cu2+ complexes; rate constants and time course for the acid-assisted dissociation of the Cu2+ complexes; voltammograms of copper(II) complexes. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] (R.D.), [email protected] (F.D.). Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors acknowledge Fundaçaõ para a Ciência e a Tecnologia (FCT), with coparticipation of the European Community funds FEDER, POCI, QREN, and COMPETE for the financial support under project PTDC/QUI/67175/ 2006, including the fellowship of C.V.E. The NMR spectrometer used is part of the National NMR Network and was purchased in the framework of the National Program for Scientific Re-equipment, contract REDE/1517/RMN/2005, with funds from POCI 2010 (FEDER) and FCT. This work was also supported by FCT through grant #PEst-OE/EQB/ LA0004/2011. Support was also provided by the CNRS, the University of Burgundy, and the Conseil Régional de Bourgogne through the 3MIM Project. P.D. thanks the French Ministry of Research for a Ph.D. grant.

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REFERENCES

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Remarkable Inertness of Copper(II) Chelates of Cyclen-Based

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